The ionization energy measures the energy changes associated with removing electrons from an atom to form positively charged ions. For example, we recall that the first ionization energy of Cl(g), 1251 kJ/mol, is the energy change associated with the following process:
The positive value of the ionization energy reminds us that energy must be put into the atom in order to remove the electron.
In addition, most atoms can gain electrons to form negatively charged ions. The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity because it measures the attraction, or affinity, of the atom for the added electron. For most atoms, energy is released when an electron is added. For example, the addition of an electron to a chlorine atom is accompanied by an energy change of -349 kJ/mol, the negative sign indicating that energy is released during the process. We therefore say that the electron affinity of Cl is -349 kJ/mol:
It is important to understand the differences between ionization energy and electron affinity: Remember that ionization energy measures the ease with which an atom loses an electron, whereas electron affinity measures the ease with which an atom gains an electron. (Two sign conventions are used for electron affinity. In most introductory texts, including this one, the thermodynamic sign convention is used: A negative sign indicates that the addition of an electron is an exothermic process, as in the electron affinity given for chlorine, -349 kJ/mol. Historically, however, electron affinity has been defined as the energy released when an electron is added to a gaseous atom or ion. Because 349 kJ/mol are released when an electron is added to Cl(g), the electron affinity by this convention is +349 kJ/mol.)
The greater the attraction between a given atom and an added electron, the more negative the atom's electron affinity will be. The electron affinity of Cl is the most negative of all the elements. For some elements, such as the noble gases, the electron affinity has a positive value, meaning that the anion is higher in energy than are the separated atom and electron:
Because E > 0, the Ar– ion is not stable and will not form.
Figure 7.8 shows the electron affinities for the representative elements in the first five rows of the periodic table. The electron affinity generally becomes increasingly negative as we proceed in each row toward the halogens. The halogens, which are one electron shy of a filled p subshell, have the most negative electron affinities. By gaining an electron, a halogen atom forms a stable negative ion that has a noble-gas configuration (Equation 7.4). The addition of an electron to a noble gas, however, would require that the electron reside in a new, higher-energy subshell (Equation 7.5). Occupying a higher-energy subshell is energetically unfavorable, so the electron affinity is positive, meaning that the ion will not form. The electron affinities of Be and Mg are positive for the same reason; the added electron would reside in a previously empty p subshell that is higher in energy.
Figure 7.8 Electron affinities in kJ/mol for the representative elements in the first five periods of the periodic table. The more negative the electron affinity, the greater the attraction of the atom for an electron. An electron affinity > 0 indicates that the negative ion is higher in energy than the separated atom and electron.
The electron affinities of the group 5A elements (N, P, As, Sb) are also interesting. Because these elements have half-filled p subshells, the added electron must be put in an orbital that is already occupied, resulting in larger electron-electron repulsions. As a result, these elements have electron affinities that are either positive (N) or less negative than their neighbors to the left (P, As, Sb).
Electron affinities do not change greatly as we move down a group. For example, consider the electron affinities of the halogens (Figure 7.8). For F, the added electron goes into a 2p orbital, for Cl a 3p orbital, for Br a 4p orbital, and so forth. Thus, as we proceed from F to I, the average distance of the added electron from the nucleus steadily increases, causing the electron-nucleus attraction to decrease. The orbital that holds the outermost electron is increasingly spread out, however, as we proceed from F to I, thereby reducing the electron-electron repulsions. A lower electron-nucleus attraction is thus counterbalanced by lower electron-electron repulsions.